# Thermochemistry - PowerPoint PPT Presentation

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Thermochemistry. Energy. Energy is necessary for all life. The study of energy and it transformations is known as thermodynamics.

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Thermochemistry

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## Thermochemistry

### Energy

• Energy is necessary for all life.

• The study of energy and it transformations is known as thermodynamics.

• In chemical reactions we study that aspect of thermodynamics that involves the relationships between chemical reactions and energy changes involving heat. This relationship is called thermochemistry.

### The nature of energy

• Energy is defined as the capacity to do work or to transfer heat.

• Work is the force required to move an object against. It is the energy required to overcome friction and inertia.

• Heat is the energy used to cause the temperature of an object to increase.

### Kinetic Energy

• Kinetic energy is the energy of motion

• Ek = ½ mv2

• The kinetic energy of an object increases as the mass increases and as the square of the velocity increases.

• A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph) has a kinetic energy of 50kJ. The same car traveling at 20 m/s (45 mph) has a kinetic energy of 200 KJ.

• ### Potential Energy

• Potential energy is the energy of position, it arises when a force operates on an object.

• The two potential energies we are familiar with are gravitational and electrostatic.

• Ep = mgh

• Where m is the mass of the object, h is the height of the object relative to a reference height and g is the gravitational constant of 9.8 m/s2

### Potential Energy (cont)

• As chemists we don’t concern ourselves with the potential energy of gravitation due to the small size of the objects concerned.

• Very concerned with electrostatic potential.

• The electrostatic potential arises from the interactions between charged particles.

• Eel = kQ1Q2/d

• Where k is a constant of proportionality = 8.99 x 109 J-m/C2

• ### Potential Energy (cont)

• The chemical energy of molecules is due to the potential energy stored in the arrangements of their atoms.

### Units of Energy

• The SI unit of energy is the joule, J

• The car from the previous example –

• A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph)

• Ek = ½ mv2

• = ½ (1000 kg)(10 m/s)2

• = ½ (1000 kg)(100 m2/s2)

• = 50,000 kg-m2/s2

• = 50 kJ

### Units of Energy

• The calorie was originally defined as the amount of energy required to raise the temperature of 1 g of water from 14.5 oC to 15.5 oC

• 1 cal = 4.184 J (exactly)

• Side note the calorie in food units is equal to 1 kcal.

### System and Surroundings

• We can not account for the entire universe when performing our analysis.

• We focus on a small portion of the universe and that portion is called the system, everything else is the surroundings.

• When we study the energy change that accompanies a chemical reaction, the reactants and products constitute the system.

### Systems

• A closed system can exchange energy but not matter with the surroundings.

• A car’s cylinder during ignition is a closed system.

• An open system can exchange energy and matter with the surroundings.

• A open beaker with boiling water

• An isolated system can exchange neither energy nor matter with the surroundings.

• An insulated thermos approximates an isolated system.

### Transferring Energy: Work & Heat

• Energy is transferred between systems and surroundings via heat or work or both.

• Energy used to cause an object to move is called work.

• W = F x d

• We perform work when we lift an object against the force of gravity.

• Heat is the energy transferred from a hotter object to a colder one. Heat always flows from hotter to colder objects.

### The First Law of Thermodynamcis

• Energy is conserved – energy can neither be created or destroyed only converted from one form to another.

• The internal energy of a system is the sum of all the kinetic energy and potential energies of all its components.

### Internal Energy

• The internal energy includes translational energies, vibrational energies, rotational energies of every atom in the system.

• The numerical value of the internal energy of a system is generally not known to a high degree of accuracy. (Why?)

• We can calculate the change in internal energy.

• ΔE = Efinal – Einitial

### Internal Energy

• ΔE contains three parts

• A number

• A unit

• A direction

• A negative value of ΔE indicates the system has lost energy to its surroundings

• A positive value of ΔE indicates the system has gained energy from its surroundings

• It is important to note that energy changes are from the point of view of the system not the surroundings.

### Internal Energy (cont)

• In a chemical reaction, the initial state of the system refers to the reactants.

• The final state of the system refers to the products.

• When hydrogen and oxygen form water at a given temperature, the system loses energy to the surroundings, thus the internal energy of the products is less than that of the reactants.

### ΔE – heat and work

• The internal energy of a system changes in magnitude as heat is added to or removed from the system or as work is done on it or by it.

• ΔE = q + w

• When heat is added to a system or work is done on a system, its internal energy increases.

### Endothermic and Exothermic Processes

• When a process occurs in which the system absorbs heat, the process is called endothermic.

• During an endothermic process heat flows into the system from its surroundings

• When a process occurs in which the system emits heat, the process is called exothermic.

• During an exothermic process heat flows from the system to its surroundings.

### State functions

• The value of a state function depends only on the present state of the system, not on the path the system took to reach that state.

• Because E is a state function, ΔE, depends only on the initial and final states of the system, not on how the change occurs.

• E is a state function however q and w are not.

### Enthalpy

• Even though most of the chemical reactions we will examine occur at a constant atmospheric pressure, work is still being performed. (w = F x d)

• As an example dissolving Zn in and acid release H2 gas. The expansion of the gas against atmospheric pressure is work.

• The work involved in the expansion or compression of a gas is called pressure-volume work.

• w = -P ΔV

### Enthalpy (cont)

• Enthalpy is a thermodynamic function that accounts for heat flow in a process occurring at constant pressure when the only work performed is P-V work.

• Enthalpy is denoted by the symbol H

• H = E + PV

• The change in enthalpy of a system equals the heat gained or lost at constant pressure.

• p 176

### Enthalpies of reactions

• Since ΔH = Hproducts – Hreactants the enthalpy change that accompanies a reaction is called the enthalpy of reaction or heat of reaction.

• When exactly 2 moles of H2 reacts with exactly 1 mole of O2 to form 2 moles of H2O at a constant pressure, the system releases 483.6 kJ of heat.

• 2H2 + O2 2H2O(g) ΔH = -483.6 kJ

• Examine energy state diagram p 177

### Thermochemical equations and enthalpy diagrams

• Enthalpy is an extensive property.

• The magnitude of ΔH is directly proportional to the amount of reactant consumed in the process.

• CH4(g) + O2(g)  CO2(g) + 2H2O(l) ΔH = -890 kJ

• If two moles of methane are consumed then

• ΔH = -1780 kJ

• The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to the ΔH of the reverse reaction.

• CO2 (g) + 2H2O(l)  CH4(g) + O2(g) ΔH = 890 kJ

### Thermochemical equations and enthalpy diagrams (cont)

• The enthalpy change for a reaction depends on the state of the reactants.

• If the products in the combustion of methane were gaseous H2O instead of liquid H2O ΔH would be -802 kJ instead of – 890 kJ.

• It is necessary to specify the state of the reactants and products.

### Calorimetry

• The value of ΔH can be determined experimentally via a calorimeter.

• The measurement of heat flow is calorimetry.

• We determine the magnitude of the heat flow by measuring the magnitude of the temperature change.

### Statements of the Obvious

• The more heat an object gains, the hotter it gets.

• All substances change temperature when they are heated.

• But less obvious is that the magnitude of the change with a given heat varies from substance to substance.

### Heat Capacity and Specific Heat

• The temperature change experienced by an object when it absorbs certain amount of heat is determined by its heat capacity, C.

• The heat capacity of an object is the amount of heat required to raise its temperature by 1oC.

• The heat capacity of a mole of a substance is it molar heat capacity. The heat capacity of one gram of a substance is its specific heat.

### Specific Heat

• Specific heat = (quantity of heat transferred)

• (gram of substance)(temp. change)

• Cs = q/(m x ΔT)

• Exercise 5.5 p 181

### Calorimetry

• Constant-Pressure calorimetry is the process of measuring the heat transfer at constant, usually atmospheric, pressure.

• Constant-Volume calorimetry is the process of measuring heat flow at constant volume in a device called a bomb calorimetery. (p 183)

### Hess’s Law

• Many enthalpies of reaction have been tabulated and from those tabulations it is possible to calculate the enthalpy of a reaction without the need for a calorimeter.

• Hess’s law states that if a reaction is carried out in a series of steps, ΔH for the overall reaction will equal the sum of the enthalpies of reaction of the individual steps.

### Enthalpies of Formation

• Enthalpy of formation is the enthalpy or heat change during formation of a compound from its constituent elements.

• The magnitude of any enthalpy change depends on the conditions of temperature, pressure and state.

• Therefore to compare enthalpies we must define a set of conditions called the standard state.

• STP – standard temperature and pressure

### Standard enthalpy change

• The standard enthalpy change of a reaction is defined as the enthalpy change when all reactants and products are in their standard states.

• By definition, the standard enthalpy of formation of the most stable form of any element is zero.

### Examples

• Example 5.10 pg 190

• Using enthalpies of formation to calculate enthalpies of reaction

### Foods

• The vast majority of the energy our bodies consume comes from carbohydrates and fats.

• C6H12O6 + 6O2 6CO2 + 6H2O ΔH = -2803 kJ

• Tristearin is a typical fat,

• 2C57H110O6(s) + 163O2(g)  114 CO2 + 110 H2O(l) ΔH = -75,520 kJ

### Fuels

• Fossil Fuels

• Natural gas

• Petroleum

• Coal